Matter and Energy
Honors Chemistry
Name______________________________
Matter
1. For each of the following, check the one box that applies.
Physical Physical Chemical Chemical
Property Change Property Change
a pond freezes over
a puddle dries up
alcohol evaporating
blue color
boiling point
burning paper
butter melting
can neutralize a base
cracking an egg
cutting paper
density
digestion of food
dissolving sugar in water
flammability
glass breaking
hardness
iron rusting
is reactive with water
leaves changing color
lighting a match
luster
magnetizing a nail
melting glass
melting point
mowing the lawn
odor
paint fading
solubility
sour taste
stretching a rubber band
supports combustion
water boiling
will form H2 when mixed with acid
wood rotting
2
Classification
2. Match each of the following statements with one or more of the following forms of matter.
E – Element
C – Compound
S – Solution
M - Mixture (heterogeneous)
_____ soil
_____ CH3OCH3
_____ perfume
_____ nitrogen gas (N2)
_____ air
_____ jello
_____ tin
_____ carbon dioxide
_____ pizza
_____ helium
_____ hot chocolate
_____ oxygen (O2)
_____ salt
_____ baking soda (NaHCO3)
_____ mercury
_____ tap water
_____ distilled water
_____ gold ring
_____ diamond
_____ sand (SiO2)
_____ liquid dish detergent
_____ calcium
_____ steam
_____ smog
_____ Italian salad dressing _____ hot chocolate, marshmallows _____ brass
Elements
3. Write the correct chemical symbol for each of the following elements.
hydrogen
tungsten
mercury
oxygen
helium
manganese
boron
sulfur
lithium
iron
aluminum
fluorine
sodium
nickel
carbon
chlorine
potassium
palladium
silicon
bromine
beryllium
platinum
tin
iodine
magnesium
copper
lead
neon
calcium
silver
nitrogen
argon
strontium
gold
phosphorus
krypton
barium
zinc
arsenic
xenon
chromium
cadmium
antimony
radon
3
Compounds
4. Complete the following table
FORMULA
a) 2NaHCO3
b)
3H3PO4
c)
(NH4)3PO4
d)
4CaCO3
e)
2Al2O3
f)
3Ba(OH)2
g)
CH3CH2OH
h)
4Ca(ClO3)2
i)
Al2(SO4)3
j)
3(NH4)2SO4
# FUNDAMENTAL
PARTICLES
2
# INDIVIDUAL ATOMS
Na: 2
H: 2
C: 2
O: 6
# TOTAL ATOMS
12
4
Mixture Separation
5. Complete the following sentences using the correct separation technique term.
filtration
crystallization (evaporation)
chromatography
distillation
a. Heterogeneous mixtures are often separated by __________________________.
b. Separating sand from water can be done by ___________________________.
c. The sugar in sugar water can be removed by ___________________________.
d. The technique that takes advantage of different boiling points is __________________.
e. Removing chlorophyll pigment from leaves might be done by ______________________.
f. Crude oil is broken down by heat, vaporized, and allowed to condense into various liquids
such as gasoline. This process is called __________________________.
6. How could you separate a mixture of iron filings and aluminum filings? What property of these
metals would allow such a separation?
7. How could you separate a mixture of alcohol, water, oil, and sand?
8. Explain what it means to say that a mixture of substances can be separated by physical means.
Give an example.
9. Explain what it means to say that a substance can be separated by chemical means. Give an
example.
5
10. A liquid sample of unknown composition arrives in your laboratory. Although the sample appears
uniform throughout, you have been told that it is a mixture of at least two different liquids. You
set up the distillation apparatus pictured below to separate the sample.
a. You notice that your sample begins to boil at 96.2°C and liquid begins to drip into the
collection flask. What is happening?
b. Why is it a good idea to remove the collection flask and replace it with another as soon as
the temperature of the solution begins to rise again?
c. The temperature rises steadily and the solution begins to boil again at 102.3°C. What is
happening now?
d. The temperature again rises until it reaches 108°C and the rest of the solution boils. How
many different liquids were contained in your sample? Explain.
e. As you complete the distillation, you notice a white, crystalline film lining the inside of the
flask that contained the original sample. What could this substance be? Why didn’t you
notice it in the original solution?
6
Percent Concentration of Solutions
11. A 197.5 g sample of a solution is found to contain 61.5 g of sodium nitrate. Calculate the percent
concentration of this solution.
12. Determine the percent concentration of a solution containing 25.0 g of table salt dissolved in 137 g
of water.
13. How many grams of water must be added to 65.5 g of glucose to make a 14.5 % solution?
14. What is the mass of water in a 23.5 % solution containing 15.9 g of potassium chlorate?
15. A student attempts to experimentally determine the percent concentration of an aqueous solution
by boiling off the water of a sample of the solution in an evaporating dish. Given the data below,
calculate the concentration of the solution.
Mass of empty evaporating dish
Mass of evaporating dish and solution
Mass of evaporating dish and dry residue
42.983 g
58.841 g
52.139 g
7
Solubility
Use the solubility curves above to answer the following questions:
a. Which salt is least soluble at 20°C?
b. How many grams of potassium chloride can be dissolved in 200g of water at 80°C?
c. At 40°C, how much potassium nitrate can be dissolved in 300 g of water?
d. Which salt shows the least change in solubility?
e. A solution at 30°C contains 90g of sodium nitrate in 100g of water. Is this solution
saturated, unsaturated, or supersaturated?
f. How many grams of potassium chlorate will crystallize when a saturated solution is cooled
from 80°C to 50°C?
g. How many grams of potassium chlorate are needed to saturate 10 g of water at 30°C?
h. Describe what will happen to an ammonia solution as it is slowly heated.
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i. Describe what will happen when 80g of potassium nitrate is added to 100g of water at
40°C and stirred.
j. A solution containing 20g of ammonium chloride dissolved in 50g of water at 100°C is
slowly cooled. At what temperature will the solution begin to crystallize?
k. List 2 chemicals from the graph that would be more soluble at higher pressures.
l. List 2 chemicals from the graph whose solubility is unaffected by pressure.
m. Describe what would happen to a small crystal of the solute that is added to a solution of
that solute that is:
i. Saturated –
ii. Unsaturated –
iii. Supersaturated –
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Energy Concepts
16. _________________ is the capacity to do work or cause change.
17. Sunlight is an example of _________________ energy.
18. The gasoline in your car is an example of ___________________ energy.
19. The SI unit for energy is the _____________.
20. The Law of Conservation of Energy states that in any process, energy is neither
________________ nor ___________________.
21. ____________________energy is the energy of motion.
22. Match the energy-requiring process with the correct energy transformation it involves.
______
______
______
______
______
______
burning gas to drive a car
friction
photosynthesis
electric mixer
steam turbine making electricity
flashlight
a. electrical to mechanical
b. light to chemical
c. mechanical to heat
d. chemical to electrical to light
e. thermal to mechanical to electrical
f. chemical to thermal to mechanical
23. If a process is endothermic, then:
a. Energy is ___________________________________.
b. The system has __________________ energy at the end of the process.
c. The temperature of the surroundings will ___________________________.
24. If a process is exothermic, then:
a. Energy is ___________________________________.
b. The system has __________________ energy at the end of the process.
c. The temperature of the surroundings will ___________________________.
.
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Phase Changes
E
VAPOR
D
CONDENSATION
C
BOILING
B
FREEZING
A
LIQUID
MELTING
SOLID
25. Answer the following questions using the heating curve above.
a. Does the temperature increase during melting?
__________________________
b. Is energy required for both melting and boiling?
__________________________
c. What state(s) of matter is(are) present at region A? __________________________
d. What state(s) of matter is(are) present at region D? __________________________
e. Is condensation endothermic or exothermic?
__________________________
f. Is melting endothermic or exothermic?
__________________________
g. How would you describe the change in arrangement of particles as heat energy and
temperature increase region A to region E.
h. What is the name given to the temperature at region D? __________________________
i. What happens to the density of water as it moves from point A to point C? _____________
j. What happens to the density of water as it moves from point C to point E? _____________
k. Describe the change in the motion of particles from regions A to E.
26. Explain the roles of intermolecular forces and kinetic energy in the states of matter and in phase
changes.
11
27. How much heat (in kJ) is required to melt a 10.0g ice cube at:
a. 0°C?
b. −15°C?
28. How many calories of heat are required to melt 157 g of gold at its melting point?
29. Aluminum has a melting point of 660.4˚C. How much heat is required to melt a 10.0 kg block of
aluminum at 77.0˚C?
30. Benzene, C6H6, boils at 80.1˚C and has a density of 0.877 g/cm3. Determine the quantity of heat,
in kilocalories, that are required to boil a 64.3 mL sample of benzene at 25.0˚C.
31. Determine the amount of heat (in kJ) required to convert 50.5 g of ice at 0°C to steam at 100°C.
12
32. How many grams of ice at 0°C can be melted by the condensation of 12.39 g of steam at 100°C,
assuming a complete transfer of thermal energy?
33. How much heat is required to convert 31.5 g of H2O from 45.0°C to 145.0°C?
34. How many grams of iron at its melting point can be melted by the addition of 39.5 kJ of heat?
35. A jet of steam at 100°C is applied to a 2.00 kg block of titanium at its melting point. How many
grams of steam would be required to completely melt the titanium assuming a complete transfer of
thermal energy?
Gas Laws
36. What happens to the volume of a gas if:
a. the pressure is doubled?
_____________________________
b. the pressure is quartered?
_____________________________
37. What happens to the pressure exerted on a gas if:
a. the volume is halved?
_____________________________
b.
_____________________________
the volume is tripled?
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38. What happens to the volume of a gas if:
a. the temperature is doubled? _____________________________
b. the temperature is quartered? _____________________________
39. What happens to the temperature of a gas if:
a. the volume is halved?
_____________________________
b. the volume is tripled?
_____________________________
40. What happens to the pressure of a gas if the temperature:
a. is quadrupled?
______________________
b. is halved?
______________________
41. What happens to the volume of a gas if:
a. the pressure is halved and the temperature is doubled?
______________________
b. the temperature is halved and the pressure is doubled?
______________________
c. both the temperature and pressure are tripled?
______________________
42. The volume of a helium balloon is 2.9 L at 91.2 kPa. Determine the volume at standard pressure.
P1 =
P2 =
V1 =
V2 =
43. A gas occupies of volume of 29.5 mL at 25°C. If the pressure remains constant, determine the
volume of the gas if the temperature is increased to 57°C.
V1 =
V2 =
T1 =
T2 =
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44. A sample of carbon dioxide has a volume of 2.88 L at 73°C and 655 torr. Determine the volume
of this gas at STP.
P1 =
P2 =
V1 =
V2 =
T1 =
T2 =
45. A gas with a pressure of 1560 torr at 198.0°C is cooled to −16.3°C. Calculate the new pressure.
46. What is the pressure on a sulfur dioxide sample with a volume of 346 mL if it occupies a volume
of 0.790 L at 125 kPa?
47. A gas has a volume of 41.3 mL at 19°C. What temperature would this gas have at 0.121 L?
48. What is the volume of a gas at STP if it occupies of volume of 14.9 L at –16.8°C and 136 kPa?
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49. The pressure of a gas at 55°C is increased from 169.2 kPa to 3.65 atm. What will its Celsius
temperature become?
50. What is the Celsius temperature of 68.2 mL of methane, if it occupies a volume of 0.0220 L at
50.0°C?
51. A sample of ammonia gas has a volume of 592 mL at 119°C and 3.71 atm. What is the Celsius
temperature if the pressure of the gas is changed to 232 kPa and its volume becomes 1.10 L?
52. A rigid container has 25.0 g of steam at 100.°C that is exerting a pressure of 93.5 kPa on the walls
of the container. If 857 J of heat are added to the container, what pressure will the steam exert?
53. A piston contains 37.9 g of steam. The piston is able to expand and contract to maintain a
constant pressure. If the steam occupies a volume of 455 mL at 100.°C, what will the volume be
after the addition of 549 cal of thermal energy?
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Dalton’s Law of Partial Pressures
54. The total pressure of a mixture of helium, neon, and hydrogen is 498 mm Hg. If the partial
pressures of helium and neon are 168 mm Hg and 192 mm Hg respectively, what is the partial
pressure of hydrogen?
55. A gas is collected over water at 20.0°C in a eudiometer tube until the pressure is equilibrated with
the atmospheric pressure. What is the pressure of the dry gas if the barometric pressure is 771torr?
56. A gas is collected over water in a eudiometer tube until the pressure is equilibrated with the
atmospheric pressure. The volume of gas in the eudiometer is 32.9 mL. The temperature of the
water is 40.0°C and the barometric pressure is 754.0 mm Hg. What is the volume of this gas at
STP?
57. A gas is collected over water in a eudiometer tube until the pressure is equilibrated with the
atmospheric pressure. The volume of gas in the eudiometer is 25.3 mL. The temperature of the
water is 50.0°C and the barometric pressure is 760.0 mm Hg. If the pressure remains constant,
what will the volume of the gas be at 75.0 °C?
17
58. Refer to the diagram below. If the valves were both opened and all three gases were allowed to
diffuse and mix equally, what would the total pressure be?
Review
59. List five properties that would be considered physical properties.
60. List three properties that would be considered chemical properties.
61. Indicate if the following changes are chemical or physical in nature.
a. Knocking down bowling pins
b. Chopping a log
c. Burning a log
d. Frying an egg
e. Water boiling
f. Dissolving sugar into water
g. Combining vinegar and baking soda
h. A precipitate is formed
i. A liquid freezes
j. A gas is produced in a reaction
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62.
Homogeneous or
Heterogeneous
Element, Compound,
Mixture, Solution
salt water
iron
salt
magnesium bromide
chocolate chip cookie
ocean water
carbon dioxide
manganese
paint
snickers bar
63. Which of the following will have a higher solubility? (Use graph on page 7)
a. Sugar and water at 50°C or at 25°C?
b. Carbon dioxide and water at 50°C or 25°C?
c. Sugar and water at 1.5 atm or 2.5 atm?
d. Carbon dioxide and water at 1.5 atm or 2.5 atm?
64. Why does a sugar cube dissolve more slowly than granulated sugar?
65. When a crystal of a solid solute is added to a solution containing that solute, the solution begins to
crystallize. What kind of solution was it before it crystallized? What kind of solution was it after
it crystallized?
66. Explain the difference between separating the components of a mixture and separating the
components of a compound. Give one concrete example of each.
67. What is the primary difference between a mixture and a solution? What do they have in common?
68. What is the primary difference between an element and a compound? What do they have in
common?
19
69. Given: 4 (NH4)3PO4
a. How many formula units are present?
__________
b. How many nitrogen atoms?
__________
c. How many hydrogen atoms?
__________
d. How many phosphorus atoms?
__________
e. How many oxygen atoms?
__________
f. How many total atoms?
__________
70. Use the given heating curve to answer the questions below
Heating Curve for Iron
4000
3500
Temperature (oC)
3000
2500
2000
1500
1000
500
0
0
2
6
4
8
10
12
Time (minutes)
a. What is the melting point for iron?
______________
b. What is the freezing point for iron?
______________
c. What is the boiling point for iron?
______________
d. What is the condensation point for iron?
______________
e. What phase is iron at room temperature?
______________
f. What phase is iron at 1000°C?
______________
g. What phase is iron at 2000°C?
______________
h. What phase is iron at 3000°C?
______________
i. What happens to the temperature of iron during melting? ___________________________
20
71. List the names of the 6 phase changes and indicate if they are endothermic or exothermic.
72. How many calories of heat are required to change 35.0 g H2O at −7.0°C to 90.0°C?
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73. Determine the volume of a gas at 25.0 atm if it has a volume of 2.30 L at 1025 mmHg.
74. The volume of a gas is 12.54 mL at 455 torr. What pressure is required to change the volume to
0.750 L?
75. A mixture carbon dioxide and carbon monoxide has a total pressure of 0.659 atm. If the carbon
dioxide has a partial pressure that is twice that of carbon monoxide, determine the partial pressure
of each gas in torr.
21
76. A gas is collected by water displacement at 23°C and the pressure is equilibrated with the air. The
eudiometer reads 45.0 mL and the barometer reads 72.3 cmHg. Convert to STP.
77. A balloon is filled to a volume of 758 mL with 169 grams of steam at 125.0°C. After a brief
period of cooling, the balloon’s volume is measured to be 745 mL. How much heat energy did the
steam lose?
Cumulative Questions
78. The density of copper was experimentally determined to be 8.52 g/mL. The accepted density of
copper is 8.92 g/mL. What is the percent error?
79. What is the density of a 55.5 g object, in g/L, if the dimensions of the object are as follows:
5.4 cm x 1.2 cm x 3.8 cm?
80. Convert the following measurements to the specified unit:
a. 2.55 nL = ____________________________ daL
b. 340 cg = ____________________________ μg
22
81. Convert 35.98 kilograms into pounds and express the answer in scientific notation.
82. A block occupies 0.2587 ft3. What is its volume in mm3?
83. If a car is traveling at 55 mph, what is its speed in nm per second?
84. A 28.9 grams sample of silicon dioxide was heated from 22.4°C to 65.0°C. How much heat
energy was absorbed by the silicon dioxide in this process?